Unit 4_ States of Matter and Gas Laws

ChemA
Unit 4: States of
Matter and Gas Laws
Chapters 12 and 13
Chapter 12
States of Matter
The Kinetic-Molecular Theory
Section 12-1
Kinetic-molecular theory explains the different properties of solids, liquids,
and gases.
• Atomic composition affects chemical properties.
• Atomic composition also affects physical properties.
• The kinetic-molecular theory describes the behavior of matter in terms of
particles in motion.
Section 12.1
The Nature of Gases
• Kinetic refers to motion
• The energy an object has because of
it’s motion is called kinetic energy
• The kinetic theory states that the tiny
particles in all forms of matter are in
constant motion!
Section 12.1
The Nature of Gases
• Three basic assumptions of the kinetic
theory as it applies to gases:
• 1. Gas is composed of particles- usually
molecules or atoms
–Small, hard spheres
–Insignificant volume; relatively far apart
from each other
–No attraction or repulsion between
particles
Section 12.1
The Nature of Gases
• 2. Particles in a gas move rapidly in
constant random motion
–Move in straight paths, changing
direction only when colliding with one
another or other objects
–Average speed of O2 in air at 20 oC is an
amazing 1660 km/h!
–Random walk is a very short distance
Section 12.1
The Nature of Gases
• 3. Collisions are perfectly elastic-
meaning kinetic energy is transferred
without loss from one particle to
another- the total kinetic energy
remains constant
The Kinetic-Molecular Theory (cont.)
Section 12-1
Gas particles are in constant random motion.
An elastic collision is one in which no kinetic energy is lost.
Section 12.1
The Nature of Gases
• Gas Pressure – defined as the force
exerted by a gas per unit surface area
of an object
–Due to: a) force of collisions, and b)
number of collisions
–No particles present? Then there cannot
be any collisions, and thus no pressure –
called a vacuum
Explaining the Behavior of Gases
Section 12-1
Great amounts of space exist between gas particles.
Compression reduces the empty spaces between particles.
Section 12.1
The Nature of Gases
• Atmospheric pressure results from the
collisions of air molecules with objects
–Decreases as you climb a mountain
because the air layer thins out as
elevation increases
• Barometer is the measuring instrument
for atmospheric pressure; dependent
upon weather
Section 12.1
The Nature of Gases
• The SI unit of pressure is the pascal
(Pa)
–At sea level, atmospheric pressure is
about 101.3 kilopascals (kPa)
–Older units of pressure include
millimeters of mercury (mm Hg), and
atmospheres (atm) – both of which came
from using a mercury barometer
Section 12.1
The Nature of Gases
• Mercury Barometer – a straight glass
tube filled with Hg, and closed at one
end; placed in a dish of Hg, with the
open end below the surface
–At sea level, the mercury would rise to
760 mm high at 25 oC- called one
standard atmosphere (atm)
Gas Pressure (cont.)
Section 12-1
Torricelli invented the barometer.
Barometers are instruments used to
measure atmospheric air
pressure.
Section 12.1
The Nature of Gases
• 1 atm = 760 mm Hg = 101.3 kPa
• Most modern barometers do not
contain mercury- too dangerous
–These are called aneroid barometers, and
contain a sensitive metal diaphragm that
responds to the number of collisions of
air molecules
Gas Pressure (cont.)
Section 12-1
Explaining the Behavior of Gases (cont.)
Section 12-1
Gases easily flow past each other because there are no significant forces of
attraction.
Diffusion is the movement of one material through another.
Effusion is a gas escaping through a tiny opening.
Diffusion
Molecules moving from areas of high
concentration to low concentration.
Example: perfume molecules
spreading across the room.
• Effusion: Gas escaping through a tiny hole in a
container.
• Depends on the speed of the molecule.
Graham’s Law
• Heavier molecules move slower at the
same temp. (by Square root)
• Heavier molecules effuse and diffuse
slower
• Helium effuses and diffuses faster than
air - escapes from balloon.
Section 12.1
The Nature of Gases
• For gases, it is important to related
measured values to standards
–Standard conditions are defined as a
temperature of 0 oC and a pressure of
101.3 kPa, or 1 atm
–This is called Standard Temperature and
Pressure, or STP
Section 12.1
The Nature of Gases
• What happens when a substance is
heated? Particles absorb energy!
–Some of the energy is stored within the
particles- this is potential energy, and
does not raise the temperature
–Remaining energy speeds up the particles
(increases average kinetic energy)- thus
increases temperature
Section 12.1
The Nature of Gases
• The particles in any collection have a
wide range of kinetic energies, from
very low to very high- but most are
somewhere in the middle, thus the
term average kinetic energy is used
–The higher the temperature, the wider
the range of kinetic energies
Section 12.1
The Nature of Gases
• An increase in the average kinetic
energy of particles causes the
temperature to rise; as it cools, the
particles tend to move more slowly,
and the average K.E. declines
–Is there a point where they slow down
enough to stop moving?
Section 12.1
The Nature of Gases
• The particles would have no kinetic
energy at that point, because they
would have no motion
–Absolute zero (0 K, or –273 oC) is the
temperature at which the motion of
particles theoretically ceases
–Never been reached, but about 0.00001 K
has been achieved
Section 12.1
The Nature of Gases
• The Kelvin temperature scale reflects a
direct relationship between
temperature and average kinetic
energy
–Particles of He gas at 200 K have twice
the average kinetic energy as particles of
He gas at 100 K
The Kinetic-Molecular Theory (cont.)
Section 12-1
Kinetic energy of a particle depends on mass and velocity.
Temperature is a measure of the average kinetic energy of the particles in a
sample of matter.
Section 12.1
The Nature of Gases
• Solids and liquids differ in their
response to temperature
–However, at any given temperature the
particles of all substances, regardless of
their physical state, have the same
average kinetic energy
Intermolecular Forces
Section 12-2
**Attractive forces between molecules cause some materials to be solids,
some to be liquids, and some to be gases at the same temperature.
Intermolecular Forces (cont.)
Section 12-2
Dispersion forces are weak forces that result from temporary shifts in density
of electrons in electron clouds.
Intermolecular Forces (cont.)
Section 12-2
Dipole-dipole forces are attractions between oppositely charged regions of
polar molecules.
Section 12-2
Intermolecular Forces (cont.)
Hydrogen bonds are special dipole-dipole attractions that occur between
molecules that contain a hydrogen atom bonded to a small, highly
electronegative atom with at least one lone pair of electrons, typically
fluorine, oxygen, or nitrogen.
Section 12-2
Intermolecular Forces (cont.)
Liquids
Section 12-3
**Forces of attraction keep molecules closely packed in a fixed volume, but
not in a fixed position.
**Liquids are much denser than gases because of the stronger intermolecular
forces holding the particles together.
Large amounts of pressure must be applied to compress liquids to very small
amounts.
Section 12.3
The Nature of Liquids
• Liquid particles are also in motion
–Liquid particles are free to slide past one
another
–Gases and liquids can both FLOW
–However, liquid particles are attracted to
each other, whereas gases are not
Section 12.3
The Nature of Liquids
• Particles of a liquid spin and vibrate
while they move, thus contributing to
their average kinetic energy
–But, most of the particles do not have
enough energy to escape into the
gaseous state; they would have to
overcome their intermolecular attractions
with other particles
Section 12.3
The Nature of Liquids
• The intermolecular attractions also
reduce the amount of space between
particles of a liquid
–Thus, liquids are more dense than gases
–Increasing pressure on liquid has hardly
an effect on it’s volume
Section 12.3
The Nature of Liquids
• Increasing the pressure also has little
effect on the volume of a liquid
–For that reason, liquids and solids are
known as the condensed states of matter
Liquids (cont.)
Section 12-3
Cohesion is the force of attraction between identical molecules.
Adhesion is the force of attraction between molecules that are different.
Capillary action is the upward movement of liquid into a narrow cylinder, or
capillary tube.
Section 12.3
The Nature of Solids
• Particles in a liquid are relatively free to
move
–Solid particles are not
• Solid particles tend to vibrate about
fixed points, rather than sliding from
place to place
Section 12.3
The Nature of Solids
• Most solids have particles packed
against one another in a highly
organized pattern
–Tend to be dense and incompressible
–Do not flow, nor take the shape of their
container
• Are still able to move, unless they
would reach absolute zero
Solids
Section 12-3
**Solids contain particles with strong attractive intermolecular forces.
Particles in a solid vibrate in a fixed position.
**Most solids are more dense than liquids.
Ice is not more dense than water.
Section 12.3
The Nature of Solids
• Some solid substances can exist in
more than one form
–Elemental carbon is an example
–1. Diamond, formed by great pressure
–2. Graphite, which is in your pencil
–3. Buckminsterfullerene (also called
“buckyballs”) arranged in hollow cages
like a soccer ball
Section 12.3
The Nature of Solids
• These are called allotropes of carbon,
because all are made of carbon, and all
are solid
• Allotropes are two or more different
molecular forms of the same element
in the same physical state
• Not all solids are crystalline, but
instead are amorphous
Section 12.3
The Nature of Solids
• Amorphous solids lack an ordered
internal structure
–Rubber, plastic, and asphalt are all
amorphous solids- their atoms are
randomly arranged
• Another example is glasses- substances
cooled to a rigid state without
crystallizing
Section 12.3
The Nature of Solids
• Glasses are sometimes called
supercooled liquids
–The irregular internal structures of
glasses are intermediate between those
of a crystalline solid and a free-flowing
liquid
–Do not melt at a definite mp, but
gradually soften when heated
Section 12.3
The Nature of Solids
• When a crystalline solid is shattered,
the fragments tend to have the same
surface angles as the original solid
• By contrast, when amorphous solids
such as glass is shattered, the
fragments have irregular angles and
jagged edges
Section 12.4
Phase Changes
• The conversion of a liquid to a gas or
vapor is called vaporization
–When this occurs at the surface of a
liquid that is not boiling, the process is
called evaporation
–Some of the particles break away and
enter the gas or vapor state; but only
those with the minimum kinetic energy
Section 12-4
Phase Changes That Require Energy (cont.)
Particles with enough energy escape from the liquid and enter the gas phase.
Section 12-4
Phase Changes That Require Energy (cont.)
Vaporization is the process by which a liquid changes to a gas or vapor.
Evaporation is vaporization only at the surface of a liquid.
Section 12.4
Phase Changes
• A liquid will also evaporate faster when
heated
–Because the added heat increases the
average kinetic energy needed to
overcome the attractive forces
–But, evaporation is a cooling process
• Cooling occurs because those with the
highest energy escape first
Section 12.4
Phase Changes
• Particles left behind have lower
average kinetic energies; thus the
temperature decreases
–Similar to removing the fastest runner
from a race- the remaining runners have
a lower average speed
• Evaporation helps to keep our skin
cooler on a hot day, unless it is very
humid on that day. Why?
Section 12.4
Phase Changes
• Evaporation of a liquid in a closed
container is somewhat different
–No particles can escape into the outside
air
–When some particles do vaporize, these
collide with the walls of the container
producing vapor pressure
Section 12.4
Phase Changes
• Eventually, some of the particles will
return to the liquid, or condense
• After a while, the number of particles
evaporating will equal the number
condensing- the space above the liquid
is now saturated with vapor
– A dynamic equilibrium exists
– Rate of evaporation = rate of condensation
Section 12.4
Phase Changes
• Note that there will still be particles
that evaporate and condense
–There will be no NET change
• An increase in temperature of a
contained liquid increases the vapor
pressure- the particles have an
increased kinetic energy, thus more
minimum energy to escape
Phase Changes That Release Energy (cont.)
Section 12-4
As energy flows from water vapor, the velocity decreases.
The process by which a gas or vapor becomes a liquid is called condensation.
.
Section 12.4
Phase Changes
• The vapor pressure of a liquid can be
determined by a device called a
manometer
• The vapor pressure of the liquid will
push the mercury into the U-tube
• A barometer is a type of manometer
Phase Changes That Require Energy (cont.)
Section 12-4
In a closed container, the pressure exerted by a vapor over a liquid is called
vapor pressure.
Section 12-4
Manometers measure gas pressure in a closed container.
Section 12.4
Phase Changes
• We now know the rate of evaporation
from an open container increases as
heat is added
–The heating allows larger numbers of
particles at the liquid’s surface to
overcome the attractive forces
–Heating allows the average kinetic energy
of all particles to increase
Section 12.4
Phase Changes
• The boiling point (bp) is the
temperature at which the vapor
pressure of the liquid is just equal to
the external pressure
–Bubbles form throughout the liquid, rise
to the surface, and escape into the air
Section 12.4
Phase Changes
• Since the boiling point is where the
vapor pressure equals external
pressure, the bp changes if the external
pressure changes
• Normal boiling point- defined as the bp
of a liquid at a pressure of 101.3 kPa
(or standard pressure)
Section 12-4
Phase Changes That Require Energy (cont.)
The boiling point is the temperature at which the vapor pressure of a liquid
equals the atmospheric pressure.
Section 12.4
Phase Changes
• Normal bp of water = 100 oC
–However, in Denver = 95 oC, since Denver
is 1600 m above sea level and average
atmospheric pressure is about 85.3 kPa
–In pressure cookers, which reduce cooking
time, water boils above 100 oC due to the
increased pressure
Section 12.4
Phase Changes
• Autoclaves, devices often used to
sterilize medical instruments, operate
much in a similar way
• Boiling is a cooling process much the
same as evaporation
–Those particles with highest KE escape
first
Section 12.4
Phase Changes
• Turning down the source of external
heat drops the liquid’s temperature
below the boiling point
• Supplying more heat allows particles to
acquire enough KE to escape- the
temperature does not go above the
boiling point, the liquid only boils faster
Section 12.4
Phase Changes
• When a solid is heated, the particles
vibrate more rapidly as the kinetic
energy increases
–The organization of particles within the
solid breaks down, and eventually the
solid melts
• The melting point (mp) is the
temperature a solid turns to liquid
Phase Changes That Require Energy (cont.)
Section 12-4
When ice is heated, the ice eventually absorbs enough energy to break the
hydrogen bonds that hold the water molecules together.
When the bonds break, the particles move apart and ice melts into water.
The melting point of a crystalline solid is the temperature at which the forces
holding the crystal lattice together are broken and it becomes a liquid.
Section 12.4
Phase Changes
• At the melting point, the disruptive
vibrations are strong enough to
overcome the interactions holding
them in a fixed position
–Melting point can be reversed by cooling
the liquid so it freezes
–Solid liquid
Section 12-4
Phase Changes That Release Energy
As heat flows from water to the surroundings, the particles lose energy.
The freezing point is the temperature at which a liquid is converted into a
crystalline solid.
A process that releases energy to the surroundings is called exothermic.
Freezing is an exothermic process.
Section 12.4
Phase Changes
• Generally, most ionic solids have high
melting points, due to the relatively strong
forces holding them together
– Sodium chloride (an ionic compound) has a
melting point = 801 oC
• Molecular compounds have relatively low
melting points **remember
molecular/covalent compounds have
weaker intermolecular forces holding
them together
Section 12.4
Phase Changes
• Hydrogen chloride (a molecular
compound) has a mp = -112 oC
• Not all solids melt- wood and cane
sugar tend to decompose when heated
• Most solid substances are crystalline in
structure
Phase Changes That Require Energy
Section 12-4
Melting occurs when heat flows into a solid object.
Heat is the transfer of energy from an object at a higher temperature to an object
at a lower temperature.
A process in which energy is absorbed from the surroundings is
called endothermic. Melting is an endothermic process.
Section 12.4
Phase Changes
• The relationship among the solid,
liquid, and vapor states (or phases) of a
substance in a sealed container are
best represented in a single graph
called a phase diagram
• Phase diagram- gives the temperature and
pressure at which a substances exists as
solid, liquid, or gas (vapor)
Section 12.4
Phase Changes
• Fig. 12.29, page 429 shows the phase
diagram for water (next slide)
–Each region represents a pure phase
–Line between regions is where the two
phases exist in equilibrium
–Triple point is where all 3 curves meet,
the conditions where all 3 phases exist in
equilibrium!
Section 12.4
Phase Changes
• With a phase diagram, the changes in
mp and bp can be determined with
changes in external pressure
• Solids, like liquids, also have a vapor
pressure
–If high enough, they can pass to a gas or
vapor without becoming a liquid
Phase Diagrams (cont.)
Section 12-4
The phase diagram for different substances are different from water.
Section 12.4
Phase Changes
• Sublimation- the change of a substance
from a solid to a vapor without passing
through the liquid state
–Examples: iodine; dry ice; moth balls;
solid air fresheners
Section 12.4
Phase Changes
• Sublimation is useful in situations such
as freeze-drying foods- such as by
freezing the freshly brewed coffee, and
then removing the water vapor by a
vacuum pump
• Also useful in separating substances-
organic chemists separate mixtures and
purify materials
Phase Changes That Release Energy (cont.)
Section 12-4
Deposition is the process by which a gas or vapor changes directly to a solid, and
is the reverse of sublimation.
Chapter 13 Gas Laws
Section 13.1
Variables that describe a Gas
• The four variables and their common
units:
1. pressure (P) in kilopascals
2. volume (V) in Liters
3. temperature (T) in Kelvin
4. number of moles (n)
82
1. Amount of Gas (number of moles)
• When we inflate a balloon, we are
adding gas molecules.
• Increasing the number of gas particles
increases the number of collisions
–thus, the pressure increases
• If temp. is constant- doubling the number of
particles doubles pressure
Pressure and the number of
molecules are directly related
• More molecules means more
collisions.
• Fewer molecules means fewer
collisions.
• Gases naturally move from areas of
high pressure to low pressure because
there is empty space to move in - spray
can is example.
Common use?
• Aerosol (spray) cans
–gas moves from higher pressure to
lower pressure
–a propellant forces the product out
–whipped cream, hair spray, paint
85
2. Volume of Gas
• In a smaller container, molecules have
less room to move.
• Hit the sides of the container more
often.
• As volume decreases, pressure
increases. (think of a syringe)
3. Temperature of Gas
• Raising the temperature of a gas
increases the pressure, if the volume is
held constant.
• The molecules hit the walls harder, and
more frequently!
• The only way to increase the
temperature at constant pressure is to
increase the volume.
Section 13.1
The Gas Laws
• These will describe HOW gases behave.
• Can be predicted by the theory.
• Amount of change can be calculated with
mathematical equations.
Section 13.1
1. Boyle’s Law
• At a constant temperature, gas pressure
and volume are inversely related.
–As one goes up the other goes down
• Formula to use: P1 x V1= P2 x V2
Boyle's Law
Section 13-1
Boyle’s law states that the volume of a fixed amount of gas held at a constant
temperature varies inversely with the pressure.
P1V1 = P2V2 where P = pressure and V = volume
Examples
• A balloon is filled with 25 L of air at 1.0
atm pressure. If the pressure is changed
to 1.5 atm what is the new volume?
• A balloon is filled with 73 L of air at 1.3
atm pressure. What pressure is needed
to change the volume to 43 L?
Section 13.1
2. Charles’s Law
• The volume of a gas is directly
proportional to the Kelvin
temperature, if the pressure is held
constant.
• Formula to use: V1/T1 = V2/T2
Charles's Law (cont.)
Section 13-1
Examples
• What is the temperature of a gas
expanded from 2.5 L at 25 ºC to 4.1L at
constant pressure?
• What is the final volume of a gas that
starts at 8.3 L and 17 ºC, and is heated
to 96 ºC?
Section 13.1
3. Gay-Lussac’s Law
• The temperature and the pressure
of a gas are directly related, at
constant volume.
• Formula to use: P1/T1 = P2/T2
Gay-Lussac's Law (cont.)
Section 13-1
Examples
• What is the pressure inside a 0.250 L can
of deodorant that starts at 25 ºC and 1.2
atm if the temperature is raised to 100
ºC?
• At what temperature will the can above
have a pressure of 2.2 atm?
Section 13.1
4. Combined Gas Law
• The Combined Gas Law deals with the
situation where only the number of
molecules stays constant.
• Formula: (P1 x V1)/T1= (P2 x V2)/T2
• This lets us figure out one thing when
two of the others change.
• The combined gas law contains all the
other gas laws!
• If the temperature remains constant...
P 1 x V1 P2 x V2
=
T1 Boyle’s Law T2
• The combined gas law contains all the
other gas laws!
• If the pressure remains constant...
P 1 x V1 P2 x V2
=
T1 T2
Charles’s Law
 The combined gas law contains
all the other gas laws!
 If the volume remains constant...
P 1 x V1 P2 x V2
=
T1 T2
Gay-Lussac’s Law
Combined Law Examples
• A 15 L cylinder of gas at 4.8 atm pressure
and 25 ºC is heated to 75 ºC and
compressed to 17 atm. What is the new
volume?
• If 6.2 L of gas at 723 mm Hg and 21 ºC is
compressed to 2.2 L at 4117 mm Hg,
what is the final temperature of the gas?
The Combined Gas Law (cont.)
Section 13-1